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Trends In the periodic table
Trends in Atomic Radii
• The atomic radius of an atom id defined as
half the distance between the nuclei of 2
atoms of the same element joined together by
a single covalent bond
Look at the atomic radii in the
following table – notice any patterns?
• In general atomic radii decrease across the
period and increase down the group
• It is important when studying any trends in the
periodic tables to remember the electrostatic
attraction between the positively charged
protons and negatively charged electrons, if the
attraction is large the protons will pull the outer
electrons nearer to the nucleus giving a smaller
atomic radius
2 Reasons for the increase in atomic
radius down a group
1. The additional electrons are going into a new
energy level thus the outer electrons are
becoming further away from the nucleus giving
a larger atomic radius
2. Screening effect of inner electrons : even
though there are more protons and electrons as
you move down a group, the inner shells of
electrons help to “shield” the outer electrons
from the charge of the nucleus allowing atomic
radii to be bigger
2 Reasons Atomic Radii decrease from
left to right across a period
1. Increasing Nuclear Charge: The number of protons in
the nucleus increase from left to right across any one
period. This has a greater attractive force on the outer
electrons drawing the energy levels closer to the
nucleus and decreasing the atomic radius
2. No increase in screening effect: Even though each
element across a period has extra electrons they are
being added to the same energy level so there is no
increase in screening effect, thus there is nothing to
“shield” the extra attractive force from the nucleus
Electronegativity
What is it?
• Electronegativity is the power of an
atom to attract electrons to itself
in a covalent bond
Electronegativity
Pauling’s electronegativity scale
• The higher the value, the more
electronegative the element
• Fluorine is the most electronegative element
• It has an electronegativity value of 4.0
Electronegativity
Pauling’s electronegativity scale
H
2.1
He
-
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Ne
-
Na
0.9
Mg
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
Ar
-
Trends in Electronegativity
• In general Electronegativity values
decrease down a group and increase
across a period
2 Reasons Electronegativity Values
Decrease down a group
1. The atomic radius increases down a group so
the outermost electrons are becoming further
away from the influence of the nucleus, this
means there is a smaller force of attraction
between the nucleus and shared pair of
electrons
2. The screening effect of inner electrons “shields”
the outermost electrons and also the shared
pair of electrons in a bond from the influence of
the nucleus. As you move down a group there is
increased screening effect
2 Reasons Electronegativity Increases
across a period
1. Increasing nuclear charge (due to more protons
in the nucleus)as you move across a period
means the attraction between the nucleus and
outermost electrons becomes greater
2. Decreasing atomic radius across a period means
the electrons in the outer most level are closer
to the nucleus and thus there is a greater
attraction between the nucleus and these
electrons
Ionisation energy
• Some elements such a sodium and potassium
lose their outermost electron very easily and
thus are very reactive
• Other elements such as gold and silver have
very little tendency to lose their electrons and
thus are very unreactive
Ionisation Energy
The first ionisation energy of an element is the
minimum energy required to remove the most
loosely bound electron from an isolated atom
of that element in its gaseous state.
Unit 7, 16
• The second ionisation energy of an element
would refer to the energy needed for the
removal of a second electron from the
positive ion
• Eg. Sodium
• First Ionisation Energy
Na
Na+ + e• Second Ionisation Energy
Na+
Na2+ + e-
Trends in Ionisation Energy
• In General Ionisation
Energy increases across a
period and decreases
down group
2 Reasons Ionisation Energy decrease
down the groups in the periodic table
1. Increasing atomic radius As the atomic radius
gets larger down a group the outermost
electrons are further away from the nucleus and
are less influenced by its attractive force thus it
is easier to remove the outermost electrons ie.
Less energy is needed
2. Screening Effect of Inner Electrons increases
down a period meaning the outermost electrons
are not as influenced by the nucleus and thus
require less energy to be removed
2 Reasons Ionisation energy Increases
across the periods in the periodic table
1. Increasing nuclear charge as the number of
protons in the nucleus increases the attraction
between the nucleus and outermost electrons is
increasing meaning more energy is required to
remove one of the electrons from the outermost
shell
2. Decreasing atomic radius As the radius of an
atom decreases the outermost electrons get
nearer the nucleus and are more influenced by
its attractive force thus more energy is required
to remove the outermost electron
Graph of the first Twenty Ionisation Energies
Unit 7, 21
Explaining the Graph
1. The maximum values are for the noble gases.
Reason:
Their atoms are very stable because of their electronic
configuration [full outer (sub) level], so it is difficult to
remove an electron.
2. The minimum values are for the group one metals (alkali metals).
Reason:
Their atoms have only one electron in their outer level, so it is
easily removed (as when this is lost it will have noble gas
configuration.) This is why group one are so reactive.
3. In general, ionisation energies increase in moving across a period from the
alkali metal to the next noble gas.
Reason:
1. Increase in nuclear charge.
(greater pull for electrons)
2. Decrease in atomic radius.
4. Ionisation energies gradually decrease in moving down a group.
Reason:
1. Increase in atomic radius.
2. Screening effect.
(This is where the inner shell or shells of electrons help to
shield the outer electrons from the positive charge in the
nucleus.
Unit 7, 22
Exceptions to Rule to 3 – Across a Period
There are two exceptions to this generalisation:
(a)
Group two elements (e.g. Be, Mg) have
abnormally high values. This is because the most
loosely bound electron comes from a full s orbital.
(e.g. 1s2, 2s2, 2p6, 3s2 in Mg) which is a relatively
stable state. When the next element in each case
(B, Al) is being ionised, the electron being removed is
the single electron in the p – orbital
(e.g. 1s2, 2s2, 2p6, 3s2, 3p1 in Al).
(b)
Group five elements also show abnormally high
values (e.g. N and P). The reason here is that the
electrons being removed are from exactly half –
filled p –orbitals, (e.g. 1s2, 2s2, 2p6, 3s2, 3p3 in P)
and the half filled orbitals are the next most stable
state after that of completely filled orbitals.
Unit 7, 23
Ionisation Energy Trends - Summary
•
Increase going across a period.
–
–
•
Decrease going down a group.
–
–
•
Increase in atomic radius.
Screening effect.
Exceptions, Group 2.
–
•
Increase in nuclear charge.
Decrease in atomic radius.
Full (outer) sublevel.
Exceptions, Group 5.
–
Half full (outer) sublevel.
Be  1s 2 2s 2
Mg  1s 2 2s 2 2p6 3s 2
Ca  1s 2 2s 2 2p6 3s 2 3p6 4s 2
N  1s 2 2s 2 2p3
P  1s 2 2s 2 2p6 3s 2 3p3
Unit 7, 24
P.T.E
Trends in
IonisationEnergy
Energies
Ionisation
Trends
Unit 7, 25
Higher Ionisation Energy Levels for the Third Period
Unit 7, 26
Reasons Ionisation energies increase
significantly for each subsequent
electron being removed from an atom
• When an ion is created the remaining electrons
are more strongly attracted to the nucleus as
ions are slightly smaller than neutral atoms
• Whenever you move to a new energy level and
try to start removing electrons from here a
considerable jump in ionisation energy is seen as
the new energy level is closer to the nucleus and
is more influenced by its attractive force (the
values will be in a similar range when removing
electrons from the same energy level)
Example 1 :
The following table gives the first ionisation energies , in KJ mol1 , of
the elements in the second period of the Periodic table.
Li
Be
B
C
N
O
F
Ne
519 900 799 1090 1400 1310 1680 2080
(i)
Explain the factors which account for the trend in ionisation energies
across a period.
(ii)
Explain why the values for boron and oxygen are exceptional.
Solution :
(i)
Increase in nuclear charge.
Decrease in atomic radius.
(ii)
The values for boron and oxygen seem exceptional as the values
for the atoms before them have abnormal values, Be due to the fact
the electron is being removed from a full s - sublevel and N as the
electron is being removed from a half full p - sublevel, both of which have
extra stability.
Unit 7, 28
Example 2 :
Explain why the first ionisation energy of oxygen atoms is greater that
that of chlorine atoms.
Solution :
Chlorine is below Oxygen in the periodic table and ionisation energies
decrease as you go downwards due to:
1. Increase in atomic radius
2. Screening effect
Example 3 :
The first ionisation energy of Sodium is 496 KJ mol1 , and the second ionsiation
energy is 4562 KJ mol1. Account for the large difference.
Solution :
1st ionisation energy: Na(1s 2 , 2s 2 , 2p6 , 3s1 )  Na  (1s 2 , 2s 2 , 2p6 )  e 
2nd ionisation energy: Na (1s 2 , 2s 2 , 2p6 )  Na2 (1s 2 , 2s 2 , 2p5 )  e 
The first ionisation energy is removing an electron from a 3s orbital after
which the electronic configuration will be that of a noble gas, with a full
outer level.
The second ionisation energy is removing an electron from a full p - orbital
in a full level which is much closer to the nucleus hence a much higher value.
Unit 7, 29
Trends in Groups
• Trends in Chemical reactivity of Alkali Metals
1. Increasing reactivity down the group as the
outer most electron is further from the
nucleus
2. Reaction with oxygen, all alkali metals react
with oxygen to form oxides
2K + ½O2
K2 O
• Lithium in air will oxidise completely to a
white lithium oxide powder in hours, it only
takes a few seconds for this to happen with
caesium!
Reaction with water
• All alkali metals react with water to form the
hydroxide of the metal and release hydrogen gas
Na + H2O
NaOH + ½H2
The more reactive the metal the more heat will be
produced and the hydrogen will catch fire this is
why you see flames ! The reaction between alkali
metals and acid is so dangerous it must never be
attempted as too much explosive hydrogen is
released
Trends in the Chemical Reactivity of
the Halogens
• The halogens are the most electronegative
elements in the periodic table, since they have
such an attraction for electrons they are not
found free in nature, Chlorine gas is made
from sodium chloride
• As fluorine is the most electronegative it is too
reactive to be kept in the school laboratory
• Reactivity of the Halogens decreases as you
move down the group
• NB Look up last two pages of this chapter in
your book!